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Lecture Presentation
Chapter 9
Chemical
Bonding I: The
Lewis Model
Sherril Soman
Grand Valley State University
© 2014 Pearson Education, Inc.
HIV-Protease
• HIV-protease is a protein synthesized by the human
immunodeficiency virus (HIV).
• This particular protein is crucial to the virus’s ability to
multiply and cause AIDS.
• Pharmaceutical companies designed molecules that would
disable HIV-protease by sticking to the molecule’s active
site—protease inhibitors.
• To design such a molecule, researchers used bonding
theories to simulate the shape of potential drug molecules
and how they would interact with the protease molecule.
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Bonding Theories
• Explain how and why atoms attach together to form
molecules
• Explain why some combinations of atoms are stable and
others are not
– Why is water H2O, not HO or H3O?
• Can be used to predict the shapes of molecules
• Can be used to predict the chemical and physical
properties of compounds
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Lewis Bonding Theory
• One of the simplest bonding theories is
called Lewis theory.
• Lewis theory emphasizes valence
electrons to explain bonding.
• Using Lewis theory, we can draw models,
called Lewis structures.
– Also known as electron dot structures
• Lewis structures allow us to predict many
properties of molecules.
– Such as molecular stability, shape, size,
and polarity
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Why Do Atoms Bond?
• Chemical bonds form because they lower the potential
energy between the charged particles that compose atoms.
• A chemical bond forms when the potential energy of the
bonded atoms is less than the potential energy of the
separate atoms.
• To calculate this potential energy, you need to consider
the following interactions:
– Nucleus-to-nucleus repulsions
– Electron-to-electron repulsions
– Nucleus-to-electron attractions
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Types of Bonds
• We can classify bonds based on the kinds of
atoms that are bonded together.
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Types of Bonding
INSERT
FIGURE
9.1
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Ionic Bonds
• When a metal atom loses electrons it becomes a cation.
– Metals have low ionization energy, making it relatively
easy to remove electrons from them.
• When a nonmetal atom gains electrons it becomes
an anion.
– Nonmetals have high electron affinities, making it
advantageous to add electrons to these atoms.
• The oppositely charged ions are then attracted to each
other, resulting in an ionic bond.
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Covalent Bonds
• Nonmetal atoms have relatively high ionization energies,
so it is difficult to remove electrons from them.
• When nonmetals bond together, it is better in terms of
potential energy for the atoms to share valence electrons.
– Potential energy is lowest when the electrons are
between the nuclei.
• Shared electrons hold the atoms together by attracting
nuclei of both atoms.
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Metallic Bonds
• The relatively low ionization energy of metals allows them
to lose electrons easily.
• The simplest theory of metallic bonding involves the metal
atoms releasing their valence electrons to be shared as a
pool by all the atoms/ions in the metal.
– An organization of metal cation islands in a sea of
electrons
– Electrons delocalized throughout the metal structure
• Bonding results from attraction of cation for the
delocalized electrons.
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Metallic Bonding
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Valence Electrons and Bonding
• Valence electrons are held most loosely.
• Chemical bonding involves the transfer or sharing of
electrons between two or more atoms.
• Because of the two previously listed facts, valence
electrons are most important in bonding.
• Lewis theory focuses on the behavior of the valence
electrons.
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Determining the Number of Valence
Electrons in an Atom
• The column number on the periodic table will tell
you how many valence electrons a main group
atom has.
– Transition elements all have two valence electrons.
Why?
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Lewis Structures of Atoms
• In a Lewis structure, we represent the valence electrons of
main-group elements as dots surrounding the symbol for
the element.
– Also known as electron dot structures
• We use the symbol of the element to represent nucleus
and inner electrons.
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Lewis Structures of Atoms
• We use dots around the symbol to represent valence
electrons.
– Pair the first two dots for the s orbital electrons.
– Put one dot on each open side for the first three p
electrons.
– Then, pair the rest of the dots for the remaining p
electrons.
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Stable Electron Arrangements
and Ion Charge
• Metals form cations by losing valence shell electrons.
• Nonmetals form anions by gaining valence electrons.
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Lewis Bonding Theory
• Atoms bond because it results in a more stable
electron configuration.
– More stable = lower potential energy
• Atoms bond together by either transferring or
sharing electrons.
• Usually this results in all atoms obtaining an outer
shell with eight electrons.
– Octet rule
– There are some exceptions to this rule—the key to
remember is to try to get an electron configuration like a
noble gas.
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Octet Rule
• When atoms bond, they tend to gain, lose, or share
electrons to result in eight valence electrons.
• ns2np6
– Noble gas configuration
• Many exceptions
– H, Li, Be, B attain an electron configuration like He.
• He = two valence electrons, a duet.
• Li loses its one valence electron.
• H shares or gains one electron.
o Though it commonly loses its one electron to become H+
• Be loses two electrons to become Be2+.
o Though it commonly shares its two electrons in covalent bonds, resulting in
four valence electrons
• B loses three electrons to become B3+.
o Though it commonly shares its three electrons in covalent bonds, resulting
in six valence electrons
– Expanded octets for elements in period 3 or below
• Using empty valence d orbitals
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Lewis Theory and Ionic Bonding
• Lewis symbols can be used to represent the
transfer of electrons from metal atom to
nonmetal atom, resulting in ions that are
attracted to each other and therefore bond.
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Lewis Theory Predictions for Ionic
Bonding
• Lewis theory predicts the number of
electrons a metal atom should lose or a
nonmetal atom should gain in order to attain
a stable electron arrangement.
– The octet rule
• This allows us to predict the formulas of
ionic compounds that result.
• It also allows us to predict the relative
strengths of the resulting ionic bonds from
Coulomb’s law.
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Energetics of Ionic Bond Formation
• The ionization energy of the metal is endothermic.
– Na(s) → Na+(g) + 1 e ─
DH° = +496 kJ/mol
• The electron affinity of the nonmetal is exothermic.
– ½Cl2(g) + 1 e ─ → Cl─(g)
DH° = −244 kJ/mol
• Generally, the ionization energy of the metal is
larger than the electron affinity of the nonmetal;
therefore, the formation of the ionic compound
should be endothermic.
• But the heat of formation of most ionic compounds
is exothermic and generally large. Why?
– Na(s) + ½Cl2(g) → NaCl(s) DH°f = −411 kJ/mol
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Ionic Bonding and the Crystal Lattice
• The extra energy that is released comes
from the formation of a structure in which
every cation is surrounded by anions, and
vice versa.
• This structure is called a crystal lattice.
• The crystal lattice is held together by the
electrostatic attraction of the cations for all
the surrounding anions.
• The crystal lattice maximizes the attractions
between cations and anions, leading to the
most stable arrangement.
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Crystal Lattice
• Electrostatic attraction is nondirectional!
– No direct anion–cation pair
• Therefore, there is no ionic molecule.
– The chemical formula is an empirical formula,
simply giving the ratio of ions based on charge
balance.
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Lattice Energy
• The extra stability that accompanies the formation of the
crystal lattice is measured as the lattice energy.
• The lattice energy is the energy released when the solid
crystal forms from separate ions in the gas state.
– Always exothermic
– Hard to measure directly, but can be calculated from
knowledge of other processes
• Lattice energy depends directly on the size of charges and
inversely on distance between ions.
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Determining Lattice Energy: The Born–
Haber Cycle
• The Born–Haber cycle is a hypothetical series of
reactions that represents the formation of an ionic
compound from its constituent elements.
• The reactions are chosen so that the change in enthalpy
of each reaction is known except for the last one, which
is the lattice energy.
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Born–Haber Cycle
• Use Hess’s law to add up enthalpy changes of other
reactions to determine the lattice energy.
 DH°f(salt) = DH°f(metal atoms, g) + DH°f(nonmetal atoms, g)
+ DH°f(cations, g) + DH°f(anions, g) + DH°(crystal lattice)
 DH°(crystal lattice) = lattice energy
– For metal atom(g)  cation(g), DH°f = first ionization
energy
• Don’t forget to add together all the ionization energies to
get to the desired cation.
o M2+ = 1st IE + 2nd IE
– For nonmetal atoms (g)  anions (g), DH°f = electron
affinity
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Born–Haber Cycle for NaCl
Na(s) → Na(g)
½ Cl2(g) → Cl(g)
Na(g) → Na+(g)
Cl (g) → Cl−(g)
Na+ (g) + Cl−(g) → NaCl(s)
Na(s) + ½ Cl2(g) → NaCl(s)
+108
DH
kJ
f(Na,g)
DHf(Cl,g)kJ)
+½(244
DHf(Na
+496
kJ+,g)
−,g)
DHf(ClkJ
−349
DH (NaCl lattice)
DHf (NaCl,
−411
kJ s)
DH°f(NaCl, s) = DH°f(Na atoms,g)
+ DH°f(Cl–Cl
(Cl atoms,
bondg)energy)
+ DH°f(Na
+ +,g)
−,g) + DH°(NaCl
+ DH°
Na
1stf(Cl
ionization
energy +lattice)Cl
electron affinity +
NaCl
lattice energy
NaCl lattice energy = (−411
DH°f(NaCl,
kJ) s) −
[DH°
−
[(+108
atoms,
+ (+122
g) + kJ) + DH°f(Cl–
f(Na kJ)
Cl bond
(+496
kJ)energy)
+ (−349
+ kJ) ] Na 1st
ionization
=
−788 kJ energy +
Cl
electron affinity ]
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Trends in Lattice Energy: Ion Size
• The force of attraction between charged particles is
inversely proportional to the distance between them.
• Larger ions mean the center of positive charge (nucleus
of the cation) is farther away from the negative charge
(electrons of the anion).
– Larger ion = weaker attraction
– Weaker attraction = smaller lattice energy
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Lattice Energy versus Ion Size
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Trends in Lattice Energy: Ion Charge
• The force of attraction between
oppositely charged particles is directly
proportional to the product of the
charges.
• Larger charge means the ions are more
strongly attracted.
– Larger charge = stronger
attraction
– Stronger attraction = larger lattice
energy
• Of the two factors, ion charge is
generally more important.
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Ionic Bonding Model versus Reality
• Lewis theory implies that the attractions between ions are
strong.
• Lewis theory predicts ionic compounds should have high
melting points and boiling points because breaking down
the crystal should require a lot of energy.
– The stronger the attraction (larger the lattice energy),
the higher the melting point.
• Ionic compounds have high melting points and boiling
points.
– MP generally > 300 °C
– All ionic compounds are solids at room temperature.
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Properties of Ionic Compounds
• Hard and brittle crystalline solids
– All are solids at room temperature
• Melting points generally > 300 C
• The liquid state conducts electricity
– The solid state does not conduct electricity
• Many are soluble in water
– The solution conducts electricity well
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Properties of Ionic Compounds
Image at top of page 390
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Ionic Bonding Model versus Reality
• Lewis theory implies that the positions of the ions in the
crystal lattice are critical to the stability of the structure.
• Lewis theory predicts that moving ions out of position should
therefore be difficult, and ionic solids should be hard.
– Hardness is measured by rubbing two materials together
and seeing which “streaks” or cuts the other.
– The harder material is the one that cuts or doesn’t streak.
• Ionic solids are relatively hard.
– Compared to most molecular solids
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Ionic Bonding Model versus Reality
• Lewis theory implies that if the ions are displaced from their
position in the crystal lattice, repulsive forces should occur.
• This predicts the crystal will become unstable and break
apart. Lewis theory predicts ionic solids will be brittle.
• Ionic solids are brittle. When struck they shatter.
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Ionic Bonding Model versus Reality
• To conduct electricity, a material must have charged
particles that are able to flow through the material.
• Lewis theory implies that, in the ionic solid, the ions are
locked in position and cannot move around.
• Lewis theory predicts that ionic solids should not conduct
electricity.
• Ionic solids do not conduct electricity.
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Ionic Bonding Model versus Reality
• Lewis theory implies that, in the liquid state or when
dissolved in water, the ions will have the ability to move
around.
• Lewis theory predicts that both a liquid ionic compound
and an ionic compound dissolved in water should conduct
electricity.
• Ionic compounds conduct electricity in the liquid state or
when dissolved in water.
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Conductivity of NaCl
In NaCl(s), the
ions are stuck in
position and not
allowed to move
to the charged
rods.
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In NaCl(aq),
the ions are
separated and
allowed to
move to the
charged rods.
Lewis Theory of Covalent Bonding
• Lewis theory implies that another way atoms can
achieve an octet of valence electrons is to share their
valence electrons with other atoms.
• The shared electrons would then count toward each
atom’s octet.
• The sharing of valence electrons is called covalent
bonding.
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Covalent Bonding: Bonding and Lone
Pair Electrons
• Electrons that are shared by atoms are called
•
bonding pairs.
Electrons that are not shared by atoms but
belong to a particular atom are called
lone pairs.
Also known as nonbonding pairs
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Single Covalent Bonds
• When two atoms share one pair of electrons, it
is called a single covalent bond.
 Two electrons
• One atom may use more than one single bond
to fulfill its octet.
 To different atoms
 H only duet
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Double Covalent Bond
• When two atoms share two pairs of
electrons the result is called a double
covalent bond.
– Four electrons
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Triple Covalent Bond
• When two atoms share three pairs of
electrons the result is called a triple
covalent bond.
– Six electrons
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Covalent Bonding: Model versus Reality
• Lewis theory implies that some combinations should be stable,
whereas others should not.
– Because the stable combinations result in “octets”
• Using these ideas of Lewis theory allows us to predict the
formulas of molecules of covalently bonded substances.
• Hydrogen and the halogens are all diatomic molecular elements,
as predicted by Lewis theory.
• Oxygen generally forms either two single bonds or a double bond
in its molecular compounds, as predicted by Lewis theory.
– Though, as we’ll see, there are some stable compounds in
which oxygen has one single bond and another where it has
a triple bond, but it still has an octet.
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Predictions of Molecular Formulas by
Lewis Theory
Oxygen is more stable when it is singly bonded to two other atoms.
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Covalent Bonding: Model versus Reality
• Lewis theory of covalent bonding implies that the
attractions between atoms are directional.
– The shared electrons are most stable between the
bonding atoms.
• Therefore, Lewis theory predicts covalently bonded
compounds will be found as individual molecules.
– Rather than an array like ionic compounds
• Compounds of nonmetals are made of individual
molecule units.
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Covalent Bonding: Model versus Reality
• Lewis theory predicts that the melting and boiling
points of molecular compounds should be
relatively low.
– This involves breaking the attractions between the
molecules, but not the bonds between the atoms.
– The covalent bonds are strong, but the attractions
between the molecules are generally weak.
• Molecular compounds have low melting points and
boiling points.
– MP generally < 300 °C
– Molecular compounds are found in all three states at
room temperature.
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Intermolecular Attractions versus Bonding
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Covalent Bonding: Model versus Reality
• Lewis theory predicts that the hardness and brittleness of
molecular compounds should vary depending on the
strength of intermolecular attractive forces.
– The kind and strength of the intermolecular attractions
varies based on many factors.
• Some molecular solids are brittle and hard, but many are
soft and waxy.
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Covalent Bonding: Model versus Reality
• Lewis theory predicts that neither molecular solids nor
liquids should conduct electricity.
– There are no charged particles around to allow the
material to conduct.
• Molecular compounds do not conduct electricity in the
solid or liquid state.
• Molecular acids conduct electricity when dissolved in
water, but not in the solid or liquid state, due to them being
ionized by the water.
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Covalent Bonding: Model versus Reality
• Lewis theory predicts that the more electrons two atoms
share, the stronger the bond should be.
• Bond strength is measured by how much energy must be
added into the bond to break it in half.
• In general, triple bonds are stronger than double bonds, and
double bonds are stronger than single bonds.
– However, Lewis theory would predict that double bonds
are twice as strong as single bonds, but the reality is they
are less than twice as strong.
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Covalent Bonding: Model versus Reality
• Lewis theory predicts that the more electrons two atoms
share, the shorter the bond should be.
– When comparing bonds to like atoms
• Bond length is determined by measuring the distance
between the nuclei of bonded atoms.
• In general, triple bonds are shorter than double bonds,
and double bonds are shorter than single bonds.
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Polar Covalent Bonding
• Covalent bonding between unlike atoms results in unequal
sharing of the electrons.
– One atom pulls the electrons in the bond closer to
its side.
– One end of the bond has larger electron density than
the other.
• The result is a polar covalent bond.
– Bond polarity
– The end with the larger electron density gets a partial
negative charge.
– The end that is electron deficient gets a partial positive
charge.
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HF

EN 2.1H
F


EN 4.0

d+ H •• F d-
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Bond Polarity
• Most bonds have some degree of sharing
and some degree of ion formation to them.
• Bonds are classified as covalent if the
amount of electron transfer is insufficient for
the material to display the classic properties
of ionic compounds.
• If the sharing is unequal enough to produce
a dipole in the bond, the bond is classified
as polar covalent.
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Electronegativity
• The ability of an atom to attract bonding
electrons to itself is called electronegativity.
• Increases across period (left to right) and
decreases down group (top to bottom)
– Fluorine is the most electronegative element.
– Francium is the least electronegative element.
– Noble gas atoms are not assigned values.
– Opposite of atomic size trend.
• The larger the difference in electronegativity,
the more polar the bond.
– Negative end toward more electronegative atom.
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Electronegativity Scale
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Electronegativity Difference and Bond Type
• If the difference in electronegativity between
bonded atoms is 0, the bond is pure covalent.
– Equal sharing
• If the difference in electronegativity between
bonded atoms is 0.1 to 0.4, the bond is
nonpolar covalent.
• If the difference in electronegativity between
bonded atoms is 0.5 to 1.9, the bond is
polar covalent.
• If difference in electronegativity between
bonded atoms is larger than or equal to 2.0, the
bond is“100%” ionic.
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Bond Polarity
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Bond Dipole Moments
• Dipole moment, m, is a measure of bond
polarity.
– A dipole is a material with a + and − end.
– it is directly proportional to the size of the partial
charges and directly proportional to the distance
between them.
 m = (q)(r)
• Not Coulomb’s law
• Measured in Debyes, D
• Generally, the more electrons two atoms
share and the larger the atoms are, the larger
the dipole moment.
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Dipole Moments
Table 9.2
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Percent Ionic Character
• The percent ionic character is the
percentage of a bond’s measured dipole
moment compared to what it would be if the
electrons were completely transferred.
• The percent ionic character indicates the
degree to which the electron is transferred.
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Lewis Structures of Molecules
• Lewis theory allows us to predict the
distribution of valence electrons in a
molecule.
• Useful for understanding the bonding in
many compounds.
• Allows us to predict shapes of molecules.
• Allows us to predict properties of molecules
and how they will interact together.
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Beware!
• Lewis theory predicts that atoms will be most stable
when they have their octet of valence electrons.
• It does not require that atoms have the same number
of lone pair electrons they had before bonding.
– First use the octet rule.
• Some atoms commonly violate the octet rule.
– Be generally has two bonds and no lone pairs in its
compounds.
– B generally has three bonds and no lone pairs in its
compounds.
– Many elements may end up with more than eight valence
electrons in their structure if they can use their empty d orbitals
for bonding.
• Expanded octet
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Lewis Structures
• Generally try to follow the common bonding
patterns.
– C = 4 bonds and 0 lone pairs, N = 3 bonds and 1 lone
pair, O= 2 bonds and 2 lone pairs, H and halogen =
1 bond,
Be = 2 bonds and 0 lone pairs, B = 3 bonds and
0 lone pairs
– Often Lewis structures with line bonds have the lone
pairs left off.
• Their presence is assumed from common bonding patterns.
• Structures that result in bonding patterns different
from the common patterns may have formal
charges.
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Formal Charge, SO2
• During bonding, atoms may end with more or fewer electrons
than the valence electrons they brought in order to fulfill
octets.
• His results in atoms having a formal charge.
FC = valence e− − nonbonding e− − ½ bonding e−
left O FC = 6 − 4 − ½ (4) = 0
S
FC = 6 − 2 − ½ (6) = +1
right O
FC = 6 − 6 − ½ (2) = −1
• Sum of all the formal charges in a molecule = 0.
– In an ion, total equals the charge
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Exceptions to the Octet Rule
• Expanded octets
– Elements with empty d orbitals can have
more than eight electrons.
• Odd number electron species (e.g., NO)
– Will have one unpaired electron
– Free-radical
– Very reactive
• Incomplete octets
– B, Al
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Resonance
• Lewis theory localizes the electrons
between the atoms that are bonding
together.
• Extensions of Lewis theory suggest that
there is some degree of delocalization of
the electrons; we call this concept
resonance.
• Delocalization of charge helps to stabilize
the molecule.
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Resonance Structures
• When there is more than one Lewis structure for
•
a molecule that differ only in the position of the
electrons, they are called resonance
structures.
The actual molecule is a combination of the
resonance forms—a resonance hybrid.
The molecule does not resonate between the two
forms, though we often draw it that way.
• Look for multiple bonds or lone pairs.
..
..
..
.. O .... S .. O ..
..
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..
..
..
.. O .. S .... O ..
..
Ozone Layer
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Resonance
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Rules of Resonance Structures
•
Resonance structures must have the same connectivity.
– Only electron positions can change.
•
Resonance structures must have the same number of
electrons.
•
Second row elements have a maximum of eight electrons.
– Bonding and nonbonding
– Third row can have expanded octet
•
Formal charges must total the same.
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Drawing Resonance Structures
1. Draw the first Lewis structure
that maximizes octets.
2. Assign formal charges.
3. Move electron pairs from atoms
with (−) formal charge toward
atoms with (+) formal charge.
4. If (+) formal charge atom
second row, only move in
electrons if you can move out
electron pairs from multiple
bond.
5. If (+) formal charge atom third
row or below, keep bringing in
electron pairs to reduce the
formal charge, even if get
expanded octet.
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−1
−1
+1
−1
−1
+1
Drawing Resonance Structures
1. Draw the first Lewis structure
that maximizes octets.
2. Assign formal charges.
3. Move electron pairs from atoms
with (−) formal charge toward
atoms with (+) formal charge.
4. If (+) formal charge atom second
row, only move in electrons if
you can move out electron pairs
from multiple bond.
5. If (+) formal charge atom third
row or below, keep bringing in
electron pairs to reduce the
formal charge, even if get
expanded octet.
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−1
+2
−1
Evaluating Resonance Structures
• Better structures have fewer formal
charges.
• Better structures have smaller formal
charges.
• Better structures have the negative
formal charge on the more
electronegative atom.
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Bond Energies
• Chemical reactions involve breaking bonds in
reactant molecules and making new bonds to
create the products.
• The DH°reaction can be estimated by comparing the
cost of breaking old bonds to the income from
making new bonds.
• The amount of energy it takes to break one mole
of a bond in a compound is called the bond
energy.
– In the gas state
– Homolytically – each atom gets ½ bonding electrons
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Trends in Bond Energies
• In general, the more electrons two atoms share, the
stronger the covalent bond.
– Must be comparing bonds between like atoms
– C≡C (837 kJ) > C═C (611 kJ) > C—C (347 kJ)
– C≡N (891 kJ) > C ═ N (615 kJ) > C—N (305 kJ)
• In general, the shorter the covalent bond, the stronger
the bond.
– Must be comparing similar types of bonds
– Br—F (237 kJ) > Br—Cl (218 kJ) > Br—Br (193 kJ)
– Bonds get weaker down the column.
– Bonds get stronger across the period.
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Using Bond Energies to Estimate DH°rxn
• The actual bond energy depends on the
surrounding atoms and other factors.
• We often use average bond energies to estimate
the DHrxn.
– Works best when all reactants and products in gas state
• Bond breaking is endothermic, DH(breaking) = +.
• Bond making is exothermic, DH(making) = −.
DHrxn = ∑ (DH(bonds broken)) + ∑ (DH(bonds formed))
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Bond Lengths
• The distance between the nuclei of
bonded atoms is called the bond length.
• Because the actual bond length depends
on the other atoms around the bond we
often use the average bond length.
– Averaged for similar bonds from many
compounds
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Trends in Bond Lengths
• In general, the more electrons two atoms share, the
shorter the covalent bond.
– Must be comparing bonds between like atoms
– C≡C (120 pm) < C═C (134 pm) < C—C (154 pm)
– C≡N (116 pm) < C═N (128 pm) < C—N (147 pm)
• Generally, bond length decreases from left to right
across period.
– C—C (154 pm) > C—N (147 pm) > C—O (143 pm)
• Generally, bond length increases down the column.
– F—F (144 pm) > Cl—Cl (198 pm) > Br—Br (228 pm)
• In general, as bonds get longer, they also get
weaker.
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Bond Lengths
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Metallic Bonds
• The low ionization energy of metals allows them to lose
electrons easily.
• The simplest theory of metallic bonding involves the metal
atoms releasing their valence electrons to be shared by all
atoms/ions in the metal.
– An organization of metal cation islands in a sea of
electrons
– Electrons delocalized throughout the metal structure
• Bonding results from attraction of the cations for the
delocalized electrons.
© 2014 Pearson Education, Inc.
Metallic Bonding
© 2014 Pearson Education, Inc.
Metallic Bonding: Model versus Reality
• This theory implies that because the electrons are
delocalized, they are able to move through the metallic
crystal.
• Because electrical conductivity takes place when
charged particles (such as electrons) are able to move
through the structure, this model predicts metallic solids
should conduct electricity well.
• Metallic solids do conduct electricity well.
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Metallic Bonding: Model versus Reality
• This theory implies heating will cause the metal ions to
vibrate faster.
• Heating will therefore make it more difficult for the
electrons to travel through the crystal.
• This theory predicts the conductivity of a metal should
decrease as its temperature increases.
• As temperature increases, the electrical conductivity of
metals decreases.
© 2014 Pearson Education, Inc.
Metallic Bonding: Model versus Reality
• Heat is a form of kinetic energy.
• Collisions between particles transfer kinetic energy from
one particle to the next.
• This model implies that the small, light electrons moving
through the metallic solid can transfer kinetic energy quicker
than larger particles locked into position, which are only able
to collide via vibrational collision.
• This model predicts metallic solids should conduct heat well.
• Metallic solids do conduct heat well.
© 2014 Pearson Education, Inc.
Metallic Bonding: Model versus Reality
• Atoms emit light when electrons jump from higher energy
levels to lower energy levels.
• This model implies that the delocalized electrons will share
a set of orbitals that belong to the entire metallic crystal.
• This model implies that the delocalized electrons on the
surface can absorb the outside light and then emit it at the
same frequency.
• This model predicts that the surface of a metallic solid
should reflect light.
• Metallic solids do reflect light.
© 2014 Pearson Education, Inc.
Metallic Bonding: Model versus Reality
• According to this model, the attractive forces that hold the
metal structure together result from the attraction of the
metal atom cores for the delocalized electrons.
• This model implies the attractive forces should not break if
positions of the atom cores shift, because the mobility of the
electrons should allow the attractions to be maintained.
• This model predicts metallic solids should be malleable and
ductile.
• Metallic solids are malleable and ductile.
© 2014 Pearson Education, Inc.
Metallic Bonding: Model versus Reality
• This model says the attractions of the core atoms for the
delocalized electrons is strong because it involves full
charges.
• In order to melt, some of the attractions holding the
metallic crystal together must be broken. In order to boil,
all the attractions must be broken.
• This model predicts that metals should have high melting
points and boiling points.
• Metals generally have high melting points and boiling
points.
– All but Hg are solids at room temperature.
© 2014 Pearson Education, Inc.
Metallic Bonding: Model versus Reality
• This model implies the attractions of the atom cores for the
delocalized electrons will be stronger when there are more
delocalized electrons.
• This model implies the attractions of the atom cores for the
delocalized electrons will be stronger when the charge on
the atom core is larger.
• This model predicts that the melting point of metals should
increase as the charge on the cation increases.
– Left to right across the period
• Melting points of metal generally increase left to right
across period.
• Na (97.72 ºC) < Mg (650 ºC) < Al (660.32 ºC)
© 2014 Pearson Education, Inc.
Metallic Bonding: Model versus Reality
• Metal ions get larger as you traverse down a
column.
• This model implies the attractions of the atom
cores for the delocalized electrons will be stronger
when the atom cores are smaller.
• This model predicts that smaller metal ions should
have higher melting points.
• This model predicts that the melting points of
metals should decrease down a column.
• Melting points of metals generally decrease down
column.
• Li (180.54 ºC) > Na (97.72 ºC) > K (63.38 ºC)
© 2014 Pearson Education, Inc.
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